Period 3: Sodium to Argon
Variations in physical properties You will recall that a large number of chemical and physical properties vary periodically with their atomic number. This idea of periodicity is the basis of the periodic table. For this section of the syllabus, we will study the trends and gradations in properties across period 3, sodium to argon. At the end of this section you will be able to relate the periodic trends in atomic properties such as ionisation energy and atomic radii to structure and bonding of the elements. The structure and bonding of the elements is in turn related to the physical properties of the elements such as melting points and densities. Periodic table But why does the atomic radius decrease as we go across the period? As we move across the period, electrons are being added to the same shell at about the same distance from the nucleus 1. Protons are also being added as we go across the period. The outer or valence electrons are pulled closer towards the nucleus as the positive charge in the nucleus increases resulting in a decrease in the atomic radius across the period. Electronegativity Electronegativity may be defined as the tendency of an atom to attract electrons Atoms may be ranked by their power to attract electrons in their chemical interactions by a numerical measure based on a scale developed by Linus Pauling (Figure 2). ''What trend is observed for electronegativities across a period? ''Notice that the electronegativities increase across period 3, from sodium to chlorine. The most electronegative elements are the non-metals, while the least electronegative elements are the metals. This was expected, because as you move across the period the effective nuclear charge increases and the atom has anincreasing ability to attract electrons. Density If you recall the density of a substance is its mass per unit volume. The atomic radius determines the volume of the elements and the crystal structure tells you how closely packed the atoms in an element are to each other. At the maximum density, the atomic radius is low and the strength of the bonding (metallic or giant covalent bond) is at a maximum. After this point, the densities begin to decrease even though the atomic radii are low and the atomic masses are higher. This decrease in density occurs because the elements on the right side of the period form simple molecular structures which are held together by weak Van der Waals forces. Electrical Conductivity Electrical conductivity increases moving from Na to a maximum at Al. They are metals This increase is due to the increasing number of electrons which can move and carry charge. The electrical conductivity then decreases sharply again as the structure of the element changes to giant covalent lattice for Si. The electrical conductivity is zero for P to Ar, which have simple molecular structures because they possess no free, delocalised electrons in their structure to carry an electric charge. Melting point The trend in the melting point across the period can be linked to the structure of the element. There is an increase in the melting points from Na to Si. Na, Mg and Al are metals and the melting points increase as the strength of the metal bonds increase due to the increased number of delocalised electrons. Si has the highest melting point of the group because of its giant covalent (molecular) structure = The first ionisation energy (in KJ/mol) is the amount of energy required to remove one mole of electrons from a mole of gaseous atoms: X(g) àX+(g) + e- As you go across period 3, from Na to Ar there is a general increase in the ionisation energy. The positive nuclear charge is increased by one unit as the atomic/proton number increases across the period. However, the electron being lost is from the same principal quantum level. The effective nuclear charge is increasing from left to right as no new quantum shell is added and therefore there is no extra shielding. Therefore the outer electron is increasingly more strongly held by the increasing positive charge of the nucleus and so, increasingly, more energy is needed remove it. There are two anomalies at Mg-Al and P-S. At these points we see a decrease in the ionisation energy. Structure, bonding and physical properties of elements in period 3 Reactions of the elements Reaction with Oxygen and Structure of Oxides The metals Na, Mg and Al burn to form a giant ionic oxide lattices sodium oxide/peroxide, magnesium oxide and aluminium oxide Na2O and Na2O2, MgO and Al2O3 respectively. Silicon Si forms a giant covalent lattice of (SiO2)n where n is very larger number Phosphorus P forms two simple molecular covalent solid oxides P4O6 and P4O10. Sulphur can form two simple molecular covalent gas molecules SO2 and SO3 Chlorine Cl forms oxide molecules of Cl2O, Cl2O7 (and others). Argon has no reaction with oxygen. The overall pattern, from left to right is: 'giant ionic lattice => giant covalent lattice > small covalent molecules'. The change in bonding character from ionic to covalent in the oxide, follows the decreasing difference in electronegativity between that of the period 3 element and oxygen. Reaction with Chlorine and Structure of Chlorides All of Na to S will combine directly on heating in chlorine to give the chloride. Sodium, magnesium and aluminium give giant ionic lattices sodium chloride Na+Cl-, magnesium chloride Mg2+(Cl-)2 and aluminium chloride Al3+(Cl-)3 respectively. Note that aluminium chloride on heating sublimes above 180oC to form small Al2Cl6 covalent dimer molecules. The non-metal elements give covalent chlorides. Silicon forms the molecular covalent liquid silicon(IV) chloride SiCl4 (silicon tetrachloride) Phosphorus forms phosphorus(III) chloride PCl3 (phosphorus trichloride) with limited chlorine and phosphorus(V) chloride PCl5 (phosphorus pentachloride) with excess chlorine. Sulphur gives disulfur dichloride S2Cl2. by direct combination (and unstable SCl2 and SCl4 can also be formed). There is no stable argon chloride. The overall pattern, from left to right across period 3 is: giant ionic lattice => polymeric covalent lattice > small covalent molecules. This is chemically characteristic of metallic > non-metallic element character. The change in bonding character from ionic to covalent in the chloride, follows the decreasing difference in electronegativity between that of the element and oxygen, as in the case of oxides. The formulae largely follow a pattern of rising formulae based on the use of all outer electrons in bonding (1-5) and then a decline in valency (oxidation state of the Period 3 element). NaCl (+1), MgCl2 (+2), AlCl3 (+3), SiCl4 (+4), PCl5 (+5), S2Cl2 (+1), Cl2, Ar no chloride So the number of atoms of chlorine combined with the Period 3 element (the valency) follows the pattern References #Hill, G., Holman, J. (2000) Chemistry in Context, 5th Ed UK: Thomas Nelson & Sons Ltd. #http://www.docbrown.info/page07/ASA2ptable5.htm